|
Modified From-http://www2.asd.k12.ak.us/hauser/curriculum/html/Chemistry/Unit%209%20Modern%20Atomic%20Theory/Handouts%20and%20Notes/Unit_09_Light_(Handout).htm
Quantized
energy
In 1900,
a German scientist named Max Planck wrote an equation to
show this the relationship between energy and frequency
of electromagnetic radiation :
E
= hn
where E
is the energy of a bit of light called a quantum,
A quantum is the smallest bit of electromagnetic
radiation that can be emitted.
It is also called a photon
of light or small “packet” of electromagnetic
radiation.
The “h” in the above equation is a very small
constant called “Planck’s constant” (6.626068 ×
10-34 J s) and “n”
is the frequency of the radiation.
Through various experiments of Planck and Albert
Einstein, it came to be accepted that light has
properties of particles as well as waves.
Planck’s “quantum” idea became the basis
for the modern understanding of atomic structure.
In the above equation, as the frequency of
radiation increases, its energy increases by the
increment “h”.
In other words, energy was not continuous, it was
quantized
– only certain energies are allowed.
Continuous energy and quantized energy can be
likened to a ramp
versus a set of stairs connecting two levels of a
building. The
ramp is analogous to continuous energy – you can sit
at any position along the ramp and thus be at any
elevation between the two levels.
The stairs are analogous to quantized or discrete
energy – you can only sit at certain elevations
between the two levels and nowhere in between.
You may sit only on the steps, not in between the
steps. Only
certain elevations are allowed.
Spectra of elements
So what
does all this stuff about waves and light and Plancks
have to do with chemistry?
It had been known for many years that when
samples of elements were heated up or energized with
electricity, they burned or glowed a certain color, not
the entire rainbow of colors like we see from white
light. (Below)
Each
element seemed to have its own characteristic color.
If we energize a sample of hydrogen with
electricity, we see a light purple color and if we
energize a sample of neon with electricity we see the
characteristic bright orange color that is so common in
“neon signs”. If
we pass this light through a prism, which separates the
colors of light like rain separates the colors of white
light into a rainbow, we see not a rainbow of continuous
colors but only certain, sharp lines of color.
We see discreet energy levels, not continuous
energy levels. Only
certain colors are seen, not the whole rainbow.
Why is this so?
Between 1911 and 1913, the Danish scientist,
Niels Bohr, tried to explain the line
spectrum of the element hydrogen which contains 5
and only 5 distinct lines of color, each with their own
energy, wavelength and frequency.
Only 3 or 4 of the lines are bright enough to
see, below. Starting
at the left, violet at 410nm, blue at 434nm, green at
486nm and red at 656nm.
Bohr
imagined hydrogen’s
lone electron as orbiting around the nucleus just like
planets orbit around the sun, but at a fixed distance
from the nucleus.
The
energy of the electron is lowest when it is quite close
to the nucleus and this state of the electron is called
the ground state
of the atom.
When an atom gains extra energy ( through heating
or electricity), the electron moves farther away from
the atom. There
is a natural attraction of the negatively charged
electron for the positively charged nucleus so it
follows that it would take energy to move the electron
away from this desirable situation.
Picture one end of a rubber band pulled snug
around a finger. The
finger is the nucleus and the other end of the rubber
band is the electron.
By stretching
the rubber band, energy is added and one end of
the band (the electron) is moved farther away from the
finger (the nucleus).
This is called an excited
state of the atom.
If one now lets go of the rubber band, it comes
slamming back into the finger and the excess energy
added a moment ago is released as heat on the finger.
The energy released depends on how much added
energy was used to stretch the rubber band in the first
place. Bohr
guessed that electrons in atoms had only certain allowed
orbits, only certain distances from the nucleus.
The electron could therefore absorb only certain
energies to take them to these fixed distances and when
this excess energy was released, only certain or
quantized energies were released as light.
This would explain why only certain colors (and
therefore certain wavelengths of light) were observed
when a hydrogen atom was excited.
The farther an electron dropped back towards the
nucleus, the more energy it would release and the
shorter the wavelength of light that would be observed.
The violet line of the hydrogen spectrum was
light of greater energy than the red line in the
spectrum because the electron had fallen much farther
than it did in another hydrogen atom that released the
red light. All
5 lines are seen since billions of atoms are being
excited to various but very limited degrees.
So, had Bohr figured it out?
Could he predict the line spectra of other
elements based on this idea of electrons jumping up and
down between only certain allowed orbits?
He could not.
He tried his theory with other atoms and it did
not work. From
this analysis, it was determined that electrons do
not move
around the nucleus in circular orbits.
In fact, it is still not known just how the
electrons move around the nucleus.
Flash
from-http://spiff.rit.edu/classes/phys301/lectures/spec_lines/Atoms_Nav.swf
Next-->Electrons as Waves
|
